Chapter XXI: Chromium, Molybdenum, Tungsten, Uranium, and Manganese

Sulphur, selenium, and tellurium belong to the uneven series of the sixth group. In the even series of this group there are known _chromium, molybdenum, tungsten, and uranium_; these give acid oxides of the type RO_{3}, like SO_{3}. Their acid properties are less sharply defined than those of sulphur, selenium, and tellurium, as is the case with all elements of the even series as compared with those of the uneven series in the same group. But still the oxides CrO_{3}, MoO_{3}, WO_{3}, and even UO_{3}, have clearly defined acid properties, and form salts of the composition MO,_n_RO_{3} with bases MO. In the case of the heavy elements, and especially of uranium, the type of oxide, UO_{3}, is less acid and more basic, because in the even series of oxides the element with the highest atomic weight always acquires a more and more pronounced basic character. Hence UO_{3} shows the properties of a base, and gives salts UO_{2}X_{2}. The basic properties of chromium, molybdenum, tungsten, and uranium are most clearly expressed in the lower oxides, which they all form. Thus chromic oxide, Cr_{2}O_{3}, is as distinct a base as alumina, Al_{2}O_{3}.

Of all these elements _chromium_ is the most widely distributed and the most frequently used. It gives chromic anhydride, CrO_{3}, and chromic oxide, Cr_{2}O_{3}--two compounds whose relative amounts of oxygen stand in the ratio 2 : 1. Chromium is, although somewhat rarely, met with in nature as a compound of one or the other type. The red chromium ore of the Urals, lead chromate or crocoisite PbCrO_{4}, was the source in which chromium was discovered by Vauquelin, who gave it this name (from the Greek word signifying colour) owing to the brilliant colours of its compounds; the chromates (salts of chromic anhydride) are red and yellow, and the chromic salts (from Cr_{2}O_{3}) green and violet. The red lead chromate is, however, a rare chromium ore found only in the Urals and in a few other localities. Chromic oxide, Cr_{2}O_{3}, is more frequently met with. In small quantities it forms the colouring matter of many minerals and rocks--for example, of some serpentines. The commonest ore, and the chief source of the chromium compounds, is the _chrome iron ore_ or chromite, which occurs in the Urals[1] and Asia Minor, California, Australia, and other localities. This is magnetic iron ore, FeO,Fe_{2}O_{3}, in which the ferric oxide is replaced by chromic oxide, its composition being FeO,Cr_{2}O_{3}. Chrome iron ore crystallises in octahedra of sp. gr. 4·4; it has a feeble metallic lustre, is of a greyish-black colour, and gives a brown powder. It is very feebly acted on by acids, but when fused with potassium acid sulphate it gives a soluble mass, which contains a chromic salt, besides potassium sulphate and ferrous sulphate. In practice the treatment of chrome iron ore is mainly carried on for the preparation of chromates, and not of chromic salts, and therefore we will trace the history of the element by beginning with chromic acid, and especially with the working up of the chrome iron ore into _potassium dichromate_, K_{2}Cr_{2}O_{7}, as the most common salt of this acid. It must be remarked that chromic anhydride, CrO_{3}, is only obtained in an anhydrous state, and is distinguished for its capacity for easily giving anhydro-salts with the alkalis, containing one, two, and even three equivalents of the anhydride to one equivalent of base. Thus among the potassium salts there is known the normal or yellow chromate, K_{2}CrO_{4}, which corresponds to, and is perfectly isomorphous with, potassium sulphate, easily forms isomorphous mixtures with it, and is not therefore suitable for a process in which it is necessary to separate the salt from a mixture containing sulphates. As in the presence of a certain excess of acid, the dichromate, K_{2}Cr_{2}O_{7} = 2K_{2}CrO_{4} + 2HX - 2KX - H_{2}O, is easily formed from K_{2}CrO_{4}, the object of the manufacturer is to produce such a dichromate, the more so as it contains a larger proportion of the elements of chromic acid than the normal salt. Finely-ground chrome iron ore, when heated with an alkali, absorbs oxygen almost as easily (Chapter III., Note 7) as a mixture of the oxides of manganese with an alkali. This absorption is due to the presence of chromic oxide, which is oxidised into the anhydride, and then combines with the alkali Cr_{2}O_{3} + O_{3} = 2CrO_{3}. As the oxidation and formation of the chromate proceeds, the mass turns _yellow_. The iron is also oxidised, but does not give ferric acid, because the capacity of the chromium for oxidation is incomparably greater than that of the iron.

[1] The working of the Ural chrome iron ore into chromium compounds has
been firmly established in Russia, thanks to the endeavours of P.
K. Ushakoff, who constructed large works for this purpose on the
river Kama, near Elabougi, where as much as 2,000 tons of ore are
treated yearly, owing to which the importation of chromium
preparations into Russia has ceased.

A mixture of lime (sometimes with potash) and chrome iron ore is heated in a reverberatory furnace, with free access of air and at a red heat for several hours, until the mass becomes yellow; it then contains normal calcium chromate, CaCr_O_{4}, which is insoluble in water in the presence of an excess of lime.[1 bis] The resultant mass is ground up, and treated with water and sulphuric acid. The excess of lime forms gypsum, and the soluble calcium dichromate, CaCr_{2}O_{7}, together with a certain amount of iron, pass into solution. The solution is poured off, and chalk added to it; this precipitates the ferric oxide (the ferrous oxide is converted into ferric oxide in the furnace) and forms a fresh quantity of gypsum, while the chromic acid remains in solution--that is, it does not form the sparingly-soluble normal salt (1 part soluble in 240 parts of water). The solution then contains a fairly pure calcium dichromate, which by double decomposition gives other chromates; for example, with a solution of potassium sulphate it gives a precipitate of calcium sulphate and a solution of potassium dichromate, which crystallises when evaporated.[2]

[1 bis] But the calcium chromate is soluble in water in the presence of
an excess of chromic acid, as may be seen from the fact that a
solution of chromic acid dissolves lime.

[2] There are many variations in the details of the manufacturing
processes, and these must be looked for in works on technical
chemistry. But we may add that the chromate may also be obtained by
slightly roasting briquettes of a mixture of chrome iron and lime,
and then leaving the resultant mass to the action of moist air
(oxygen is absorbed, and the mass turns yellow).

_Potassium dichromate_, K_{2}Cr_{2}O_{7}, easily crystallises from acid solutions in red, well-formed prismatic crystals, which fuse at a red heat and evolve oxygen at a very high temperature, leaving chromic oxide and the normal salt, which undergoes no further change: 2K_{2}Cr_{2}O_{7} = 2K_{2}CrO_{4} + Cr_{2}O_{3} + O_{3}. At the ordinary temperature 100 parts of water dissolve 10 parts of this salt, and the solubility increases as the temperature rises. It is most important to note that the dichromate does not contain water, it is K_{2}CrO_{4} + CrO_{3}; the acid salt corresponding to potassium acid sulphate, KHSO_{4}, does not exist. It does not even evolve heat when dissolving in water, but on the contrary produces cold, _i.e._ it does not form a very stable compound with water. The solution and the salt itself are poisonous, and act as powerful oxidising agents, which is the character of chromic acid in general. When heated with sulphur or organic substances, with sulphurous anhydride, hydrogen sulphide, &c., this salt is deoxidised, yielding chromic compounds.[2 bis] Potassium dichromate[3] is used in the arts and in chemistry as a source for the preparation of all other chromium compounds. It is converted into yellow pigments by means of double decomposition with salts of lead, barium, and zinc. When solutions of the salts of these metals are mixed with potassium dichromate (in dyeing generally mixed with soda, in order to obtain normal salts), they are precipitated as insoluble normal salts; for example, 2BaCl_{2} + K_{2}Cr_{2}O_{7} + H_{2}O = 2BaCrO_{4} + 2KCl + 2HCl. It follows from this that these salts are insoluble in dilute acids, but the precipitation is not complete (as it would be with the normal salt). The barium and zinc salts are of a lemon yellow colour; the lead salt has a still more intense colour passing into orange. Yellow cotton prints are dyed with this pigment. The silver salt, Ag_{2}CrO_{4}, is of a bright red colour.

[2 bis] The oxidising action of potassium dichromate on organic
substances at the ordinary temperature is especially marked under
the action of light. Thus it acts on gelatin, as Poutven
discovered; this is applied to photography in the processes of
photogravure, photo-lithography, pigment printing, &c. Under the
action of light this gelatin is oxidised, and the chromic anhydride
deoxidised into chromic oxide, which unites with the gelatin and
forms a compound insoluble in warm water, whilst where the light
has not acted, the gelatin remains soluble, its properties being
unaffected by the presence of chromic acid or potassium dichromate.

[3] Ammonium and sodium dichromates are now also prepared on a large
scale. The sodium salts may be prepared in exactly the same manner
as those of potassium. The normal salt combines with ten
equivalents of water, like Glauber's salt, with which it is
isomorphous. Its solution above 30° deposits the anhydrous salt.
Sodium dichromate crystals contain Na_{2}Cr_{2}O_{7},2H_{2}O. The
_ammonium salts of chromic acid_ are obtained by saturating the
anhydride itself with ammonia. The dichromate is obtained by
saturating one part of the anhydride with ammonia, and then adding
a second part of anhydride and evaporating under the receiver of an
air-pump. On ignition, the normal and acid salts leave chromic
oxide. Potassium ammonium chromate, NH_{4}KCrO_{4}, is obtained in
yellow needles from a solution of potassium dichromate in aqueous
ammonia; it not only loses ammonia and becomes converted into
potassium dichromate when ignited, but also by degrees at the
ordinary temperature. This shows the feeble energy of chromic acid,
and its tendency to form stable dichromates. Magnesium chromate is
soluble in water, as also is the strontium salt. The calcium salt
is also somewhat soluble, but the barium salt is almost insoluble.
The isomorphism with sulphuric acid is shown in the chromates by
the fact that the magnesium and ammonium salts form double salts
containing six equivalents of water, which are perfectly
isomorphous with the corresponding sulphates. The magnesium salt
crystallises in large crystals containing seven equivalents of
water. The beryllium, cerium, and cobalt salts are insoluble in
water. Chromic acid dissolves manganous carbonate, but on
evaporation the solution deposits manganese dioxide, formed at the
expense of the oxygen of the chromic acid. Chromic acid also
oxidises ferrous oxide, and ferric oxide is soluble in chromic
acid.

One of the chromates most used by the dyer is the insoluble yellow
lead chromate, PbCrO_{4} (Chapter XVIII., Note 46), which is
precipitated on mixing solutions of PbX_{2} with soluble chromates.
It easily forms a basic salt, having the composition PbO,PbCrO_{4},
as a crystalline powder, obtained by fusing the normal salt with
nitre and then rapidly washing in water. The same substance is
obtained, although impure and in small quantity, by treating lead
chromate with neutral potassium chromate, especially on boiling the
mixture; and this gives the possibility of attaining, by means of
these materials, various tints of lead chromate, from yellow to
red, passing through different orange shades. The decomposition
which takes place (incompletely) in this case is as follows:
2PbCrO_{4} + K_{2}CrO_{4} = PbCrO_{4},PbO + K_{2}Cr_{2}O_{7}--that
is, potassium dichromate is formed in solution.

When potassium dichromate is mixed with potassium hydroxide or carbonate (carbonic anhydride being disengaged in the latter case) it forms the _normal_ salt, K_{2}CrO_{4}, known as _yellow chromate of potassium_. Its specific gravity is 2·7, being almost the same as that of the dichromate. It absorbs heat in dissolving; one part of the salt dissolves in 1·75 part of water at the ordinary temperature, forming a yellow solution. When mixed even with such feeble acids as acetic, and more especially with the ordinary acids, it gives the dichromate, and Graham obtained a trichromate, K_{2}Cr_{3}O_{10} = K_{2}CrO_{4},2CrO_{3}, by mixing a solution of the latter salt with an excess of nitric acid.

_Chromic anhydride_ is obtained by preparing a saturated solution of potassium dichromate at the ordinary temperature, and pouring it in a thin stream into an equal volume of pure sulphuric acid.[4] On mixing, the temperature naturally rises; when slowly cooled, the solution deposits chromic anhydride in needle-shaped crystals of a red colour sometimes several centimetres long. The crystals are freed from the mother liquor by placing them on a porous tile.[4 bis] It is very important at this point to call attention to the fact that a hydrate of chromic anhydride is never obtained in the decomposition of chromic compounds, but always the _anhydride_, CrO_{3}. The corresponding hydrate, CrO_{4}H_{2}, or any other hydrate, is not even known. Nevertheless, it must be admitted that chromic acid is bibasic, because it forms salts isomorphous or perfectly analogous with the salts formed by sulphuric acid, which is the best example of a bibasic acid. A clear proof of the bibasicity of CrO_{3} is seen in the fact that the anhydride and salts give (when heated with sodium chloride and sulphuric acid) a volatile chloranhydride, CrO_{2}Cl_{2}, containing two atoms of chlorine as a bibasic acid should.[5] Chromic anhydride is a red crystalline substance, which is converted into a black mass by heat; it fuses at 190°, and disengages oxygen above 250°, leaving a residue of chromium dioxide, CrO_{2},[6] and, on still further heating, chromic oxide, Cr_{2}O_{3}. Chromic anhydride is exceedingly soluble in water, and even attracts moisture from the air, but, as was mentioned above, it does not form any definite compound with water. The specific gravity of its crystals is 2·7, and when fused it has a specific gravity 2·6. The solution presents perfectly defined acid properties. It liberates carbonic anhydride from carbonates; gives insoluble precipitates of the chromates with salts of barium, lead, silver, and mercury.

[4] The sulphuric acid should not contain any lower oxides of nitrogen,
because they reduce chromic anhydride into chromic oxide. If a
solution of a chromate be heated with an excess of acid--for
instance, sulphuric or hydrochloric acid--oxygen or chlorine is
evolved, and a solution of a chromic salt is formed. Hence, under
these circumstances, chromic acid cannot be obtained from its
salts. One of the first methods employed consisted in converting
its salts into volatile _chromium hexafluoride_, CrF_{6}. This
compound, obtained by Unverdorben, may be prepared by mixing lead
chromate with fluor spar in a dry state, and treating the mixture
with fuming sulphuric acid in a platinum vessel: PbCrO_{4} +
3CaF_{2} + 4H_{2}SO_{4} = PbSO_{4} + 3CaSO_{4} + 4H_{2}O + CrF_{6}.
Fuming sulphuric acid is taken, and in considerable excess, because
the chromium fluoride which is formed is very easily decomposed by
water. It is volatile, and forms a very caustic, poisonous vapour,
which condenses when cooled in a dry platinum vessel into a red,
exceedingly volatile liquid, which fumes powerfully in air. The
vapours of this substance when introduced into water are decomposed
into hydrofluoric acid and chromic anhydride: CrF_{6} + 3H_{2}O =
CrO_{3} + 6HF. If very little water be taken the hydrofluoric acid
volatilises, and chromic anhydride separates directly in crystals.
The chloranhydride of chromic acid, CrO_{2}Cl_{2} (Note 5), is also
decomposed in the same manner. A solution of chromic acid and a
precipitate of barium sulphate are formed by treating the insoluble
barium chromate with an equivalent quantity of sulphuric acid. If
carefully evaporated, the solution yields crystals of chromic
anhydride. Fritzsche gave a very convenient method of preparing
chromic anhydride, based on the relation of chromic to sulphuric
acid. At the ordinary temperature the strong acid dissolves both
chromic anhydride and potassium chromate, but if a certain amount
of water is added to the solution the chromic anhydride separates,
and if the amount of water be increased the precipitated chromic
anhydride is again dissolved. The chromic anhydride is almost all
separated from the solution when it contains two equivalents of
water to one equivalent of sulphuric acid. Many methods for the
preparation of chromic anhydride are based on this fact.

[4 bis] They cannot be filtered through paper or washed, because the
chromic anhydride is reduced by the filter-paper, and is dissolved
during the process of washing.

[5] Berzelius observed, and Rose carefully investigated, this
remarkable reaction, which occurs between chromic acid and sodium
chloride in the presence of sulphuric acid. If 10 parts of common
salt be mixed with 12 parts of potassium dichromate, fused, cooled,
and broken up into lumps, and placed in a retort with 20 parts of
fuming sulphuric acid, it gives rise to a violent reaction,
accompanied by the formation of brown fumes of _chromic
chloranhydride_, or _chromyl chloride_, CrO_{2}Cl_{2}, according to
the reaction: CrO_{3} + 2NaCl + H_{2}SO_{4} = Na_{2}SO_{4} + H_{2}O
+ CrO_{2}Cl_{2}. The addition of an excess of sulphuric acid is
necessary in order to retain the water. The same substance is
always formed when a metallic chloride is heated with chromic acid,
or any of its salts, in the presence of sulphuric acid. The
formation of this volatile substance is easily observed from the
brown colour which is proper to its vapour. On condensing the
vapour in a dry receiver a liquid is obtained having a sp. gr. of
1·9, boiling at 118°, and giving a vapour whose density, compared
with hydrogen, is 78, which corresponds with the above formula.
Chromyl chloride is decomposed by heat into chromic oxide, oxygen,
and chlorine: 2CrO_{2}Cl_{2} = Cr_{2}O_{3} + 2Cl_{2} + O; so that
it is able to act simultaneously as a powerful oxidising and
chlorinating agent, which is taken advantage of in the
investigation of many, and especially of organic, substances. When
treated with water, this substance first falls to the bottom, and
is then decomposed into hydrochloric and chromic acids, like all
chloranhydrides: CrO_{2}Cl_{2} + H_{2}O = CrO_{3} + 2HCl. When
brought into contact with inflammable substances it sets fire to
them; it acts thus, for instance, on phosphorus, sulphur, oil of
turpentine, ammonia, hydrogen, and other substances. It attracts
moisture from the atmosphere with great energy, and must therefore
be kept in closed vessels. It dissolves iodine and chlorine, and
even forms a solid compound with the latter, which depends upon the
faculty of chromium to form its higher oxide, Cr_{2}O_{7}. The
close analogy in the physical properties of the chloranhydrides,
CrO_{2}Cl_{2} and SO_{2}Cl_{2}, is very remarkable, although
sulphurous anhydride is a gas, and the corresponding oxide,
CrO_{2}, is a non-volatile solid. It may be imagined, therefore,
that chromium dioxide (which will be mentioned in the following
note) presents a polymerised modification of the substance having
the composition CrO_{2}; in fact, this is obvious from the method
of its formation.

If three parts of potassium dichromate be mixed with four parts of
strong hydrochloric acid and a small quantity of water, and gently
warmed, it all passes into solution, and no chlorine is evolved; on
cooling, the liquid deposits red prismatic crystals, known as
_Peligot's salt_, very stable in air. This has the composition
KCl,CrO_{3}, and is formed according to the equation
K_{2}Cr_{2}O_{7} + 2HCl = 2KCl,CrO_{3} + H_{2}O. It is evident that
this is the first chloranhydride of chromic acid, HCrO_{3}Cl, in
which the hydrogen is replaced by potassium. It is decomposed by
water, and on evaporation the solution yields potassium dichromate
and hydrochloric acid. This is a fresh instance of the reversible
reactions so frequently encountered. With sulphuric acid Peligot's
salt forms chromyl chloride. The latter circumstance, and the fact
that Geuther produced Peligot's salt from potassium chromate and
chromyl chloride, give reason for thinking that it is a compound of
these two substances: 2KCl,CrO_{3} = K_{2}CrO_{4} + CrO_{2}Cl_{2}.
It is also sometimes regarded as potassium dichromate in which one
atom of oxygen is replaced by chlorine--that is,
K_{2}Cr_{2}O_{6}Cl_{2}, corresponding with K_{2}Cr_{2}O_{7}. When
heated it parts with all its chlorine, and on further heating gives
chromic oxide.

[6] This intermediate degree of oxidation, CrO_{2}, may also be
obtained by mixing solutions of chromic salts with solutions of
chromates. The brown precipitate formed contains a compound,
Cr_{2}O_{3},CrO_{3}, consisting of equivalent amounts of chromic
oxide and anhydride. The brown precipitate of chromium dioxide
contains water. The same substance is formed by the imperfect
deoxidation of chromic anhydride by various reducing agents.
Chromic oxide, when heated, absorbs oxygen, and appears to give the
same substance. Chromic nitrate, when ignited, also gives this
substance. When this substance is heated it first disengages water
and then oxygen, chromic oxide being left. It corresponds with
manganese dioxide, Cr_{2}O_{3},CrO_{3} = 3CrO_{2}. Krüger treated
chromium dioxide with a mixture of sodium chloride and sulphuric
acid, and found that chlorine gas was evolved, but that chromyl
chloride was not formed. Under the action of light, a solution of
chromic acid also deposits the brown dioxide. At the ordinary
temperature chromic anhydride leaves a brown stain upon the skin
and tissues, which probably proceeds from a decomposition of the
same kind. Chromic anhydride is soluble in alcohol containing
water, and this solution is decomposed in a similar manner by
light. Chromium dioxide forms K_{2}CrO_{4} when treated with
H_{2}O_{2} in the presence of KHO.

The action of hydrogen peroxide on a solution of chromic acid or of potassium dichromate gives a blue solution, which very quickly becomes colourless with the disengagement of oxygen. Barreswill showed that this is due to the formation of a _perchromic anhydride_, Cr_{2}O_{7}, corresponding with sulphur peroxide. This peroxide is remarkable from the fact that it very easily dissolves in ether and is much more stable in this solution, so that, by shaking up hydrogen peroxide mixed with a small quantity of chromic acid, with ether, it is possible to transfer all the blue substance formed to the ether.[6 bis]

[6 bis] Now that persulphuric acid H_{2}S_{2}O_{8} is well known it
might be supposed that perchromic anhydride, Cr_{2}O_{7}, would
correspond to perchromic acid, H_{2}Cr_{2}O_{8}, but as yet it is
not certain whether corresponding salts are formed. Péchard (1891)
on adding an excess of H_{2}O_{2} and baryta water to a dilute
solution of CrO_{2} (8 grm. per litre), observed the formation of a
yellow precipitate, but oxygen was disengaged at the same time and
the precipitate (which easily exploded when dried) was found to
contain, besides an admixture of BaO_{2}, a compound BaCrO_{5}, and
this = BaO_{2} + CrO_{3}, and does not correspond to perchromic
acid. The fact of its decomposing with an explosion, and the mode
of its preparation, proves, however, that this is a similar
derivative of peroxide of hydrogen to persulphuric acid (Chapter
XX.)

With oxygen acids, chromic acid evolves oxygen; for example, with sulphuric acid the following reaction takes place: 2CrO_{3} + 3H_{2}SO_{4} = Cr_{2}(SO_{4})_{3} + O_{3} + 3H_{2}O. It will be readily understood from this that _a mixture of chromic acid_ or _of its salts with sulphuric acid_ forms an excellent _oxidising agent_, which is frequently employed in chemical laboratories and even for technical purposes as a means of oxidation. Thus hydrogen sulphide and sulphurous anhydride are converted into sulphuric acid by this means. Chromic acid is able to act as a powerful oxidising agent because it passes into chromic oxide, and in so doing disengages half of the oxygen contained in it: 2CrO_{3} = Cr_{2}O_{3} + O_{3}. Thus chromic anhydride itself is a powerful oxidising agent, and is therefore employed instead of nitric acid in galvanic batteries (as a depolariser), the hydrogen evolved at the carbon being then oxidised, and the chromic acid converted into a non-volatile product of deoxidation, instead of yielding, as nitric acid does, volatile lower oxides of offensive odour. Organic substances are more or less perfectly oxidised by means of chromic anhydride, although this generally requires the aid of heat, and does not proceed in the presence of alkalis, but generally _in the presence of acids_. In acting on a solution of potassium iodide, chromic acid, like many oxidising agents, liberates iodine; the reaction proceeds in proportion to the amount of CrO_{3} present, and may serve for determining the amount of CrO_{3}, since the amount of iodine liberated can be accurately determined by the iodometric method (Chapter XX., Note 42). If chromic anhydride be ignited in a stream of ammonia, it gives chromic oxide, water, and nitrogen. In all cases when chromic acid acts as an oxidising agent in the presence of acids and under the action of heat, the product of its deoxidation is a chromic salt, CrX_{3}, which is characterised by the green colour of its solution, so that the _red_ or yellow _solution_ of a salt of chromic acid is then transformed into a _green solution_ of a chromic salt, derived from chromic oxide, Cr_{2}O_{3}, which is closely analogous to Al_{2}O_{3}, Fe_{2}O_{3}, and other bases of the composition R_{2}O_{3}. This analogy is seen in the insolubility of the anhydrous oxide, in the gelatinous form of the colloidal hydrate, in the formation of alums,[7] of a volatile chloride of chromium, &c.[7 bis]

[7] As a mixture of potassium dichromate and sulphuric acid is usually
employed for oxidation, the resultant solution generally contains a
double sulphate of potassium and chromium--that is, _chrome alum_,
isomorphous with ordinary alum--K_{2}Cr_{2}O_{7} + 4H_{2}SO_{4} +
20H_{2}O = O_{3} + K_{2}Cr_{2}(SO_{4})_{4},24H_{2}O or
2(KCr(SO_{4})_{2},12H_{2}O). It is prepared by dissolving potassium
dichromate in dilute sulphuric acid; alcohol is then added and the
solution slightly heated, or sulphurous anhydride is passed through
it. On the addition of alcohol to a cold mixture of potassium
dichromate and sulphuric acid, the gradual disengagement of
pleasant-smelling volatile products of the oxidation of alcohol,
and especially of aldehyde, C_{2}H_{4}O, is remarked. If the
temperature of decomposition does not exceed 35°, a _violet_
solution of chrome alum is obtained, but if the temperature be
higher, a solution of the same alum is obtained of a _green_
colour. As chrome alum requires for solution 7 parts of water at
the ordinary temperature, it follows that if a somewhat strong
solution of potassium dichromate be taken (4 parts of water and
1-1/2 of sulphuric acid to 1 part of dichromate), it will give so
concentrated a solution of chrome alum that on cooling, the salt
will separate without further evaporation. _If the liquid_,
prepared as above or in any instance of the deoxidation of chromic
acid, _be heated_ (the oxidation naturally proceeds more rapidly)
somewhat strongly, for instance, to the boiling-point of water, or
if the violet solution already formed be raised to the same
temperature, it acquires a bright _green colour_, and on
evaporation the same mixture, which at lower temperatures so easily
gives cubical crystals of chrome alum, _does not give any crystals
whatever_. _If the green solution be kept_, however, _for several
weeks_ at the ordinary temperature, it deposits _violet crystals_
of chrome alum. The green solution, when evaporated, gives a
non-crystalline mass, and the violet crystals lose water at 100°
and turn green. It must be remarked that the transition of the
green modification into the violet is accompanied by a decrease in
volume (Lecoq de Boisbaudran, Favre). If the green mass formed at
the higher temperature be evaporated to dryness and heated at 30°
in a current of air, it does not retain more then 6 equivalents of
water. Hence Löwel, and also Schrötter, concluded that the green
and violet modifications of the alum depend on different degrees of
combination with water, which may be likened to the different
compounds of sodium sulphate with water and to the different
hydrates of ferric oxide.

However, the question in this case is not so simple, as we shall
afterwards see. Not chrome alum alone, but _all the chromic salts_,
give two, if not three, _varieties_. At least, there is no doubt
about the existence of two--a _green_ and a _violet modification_.
The green chromic salts are obtained by heating solutions of the
violet salts, the violet solutions are produced on keeping
solutions of the green salts for a long time. The conversion of the
violet salts into green by the action of heat is itself an
indication of the possibility of explaining the different
modifications by their containing different proportions (or states)
of water, and, moreover, by the green salts having a less amount of
water than the violet. However, there are other explanations.
Chromic oxide is a base like alumina, and is therefore able to give
both acid and basic salts. It is supposed that the difference
between the green and violet salts is due to this fact. This
opinion of Krüger is based on the fact that alcohol separates out a
salt from the green solution which contains less sulphuric acid
than the normal violet salt. On the other hand, Löwel showed that
all the acid cannot be separated from the green chromic salts by
suitable reagents, as easily as it can be from the same solution of
the violet salts; thus barium salts do not precipitate all the
sulphuric acid from solutions of the green salts. According to
other researches the cause of the varieties of the chromic salts
lies in a difference in the bases they contain--that is, it is
connected with a modification of the properties of the oxide of
chromium itself. This only refers to the hydroxides, but as
hydroxides themselves are only special forms of salts, the
differences observed as yet in this direction between the
hydroxides only confirm the generality of the difference observed
in the chromic compounds (_see_ Note 7 bis).

The salts of chromic oxide, like those of alumina, are easily
decomposed, give basic and double salts, and have an acid reaction,
as chromic oxide is a feeble base. Potassium and sodium hydroxides
give a _precipitate_ of the hydroxide with chromic salts, CrX_{3}.
The violet and green salts give a _hydroxide soluble in an excess
of the reagent_; but the hydroxide is held in solution by very
feeble affinities, so that it is partially separated by heat and
dilution with water, and completely so on boiling. In an alkaline
solution, chromic hydroxide is easily converted into chromic acid
by the action of lead dioxide, chlorine, and other oxidising
agents. If the chromic oxide occurs together with such oxides as
magnesia, or zinc oxide, then on precipitation it separates out
from its solution in combination with these oxides, forming, for
example, ZnO,Cr_{2}O_{3}. Viard obtained compounds of Cr_{2}O_{3}
with the oxides of Mg, Zn, Cd, &c.) On precipitating the violet
solution of chrome alum with ammonia, a precipitate containing
Cr_{2}O_{3},6H_{2}O is obtained, whilst the precipitate from the
boiling solution with caustic potash was a hydrate containing four
equivalents of water. When fused with borax chromic salts give a
green glass. The same coloration is communicated to ordinary glass
by the presence of traces of chromic oxide. A chrome glass
containing a large amount of chromic oxide may be ground up and
used as a green pigment. Among the hydrates of oxide of chromium
_Guignet's green_ forms one of the widely-used green pigments which
have been substituted for the poisonous arsenical copper pigments,
such as Schweinfurt green, which formerly was much used. Guignet's
green has an extremely bright green colour, and is distinguished
for its great stability, not only under the action of light but
also towards reagents; thus it is not altered by alkaline
solutions, and even nitric acid does not act on it. This pigment
remains unchanged up to a temperature of 250°; it contains
Cr_{2}O_{3},2H_{2}O, and generally a small amount of alkali. It is
prepared by fusing 3 parts of boric acid with 1 part of potassium
dichromate; oxygen is disengaged, and a green glass, containing a
mixture of the borates of chromium and potassium, is obtained. When
cool this glass is ground up and treated with water, which extracts
the boric acid and alkali and leaves the above-named chromic
hydroxide behind. This hydroxide only parts with its water at a red
heat, leaving the anhydrous oxide.

The chromic hydroxides lose their water by ignition, and in so
doing become spontaneously incandescent, like the ordinary ferric
hydroxide (Chapter XXII.). It is not known, however, whether all
the modifications of chromic oxide show this phenomenon. The
anhydrous _chromic oxide_, Cr_{2}O_{3}, is exceedingly difficultly
soluble in acids, if it has passed through the above recalescence.
But if it has parted with its water, or the greater part of it, and
not yet undergone this self-induced incandescence (has not lost a
portion of its energy), then it is soluble in acids. It is not
reduced by hydrogen. It is easily obtained in various crystalline
forms by many methods. The chromates of mercury and ammonium give a
very convenient method for its preparation, because when ignited
they leave chromic oxide behind. In the first instance oxygen and
mercury are disengaged, and in the second case nitrogen and water:
2Hg_{2}CrO_{4} = Cr_{2}O_{3} + O_{5} + 4Hg or
(NH_{4})_{2}Cr_{2}O_{7} = Cr_{2}O_{3} + 4H_{2}O + N_{2}. The second
reaction is very energetic, and the mass of salt burns
spontaneously if the temperature be sufficiently high. A mixture of
potassium sulphate and chromic oxide is formed by heating potassium
dichromate with an equal weight of sulphur: K_{2}Cr_{2}O_{7} + S =
K_{2}SO_{4} + Cr_{2}O_{3}. The sulphate is easily extracted by
water, and there remains a bright green residue of the oxide, whose
colour is more brilliant the lower the temperature of the
decomposition. The oxide thus obtained is used as a green pigment
for china and enamel. The anhydrous chromic oxide obtained from
chromyl chloride, CrO_{2}Cl_{2}, has a specific gravity of 5·21,
and forms almost black crystals, which give a green powder. They
are hard enough to scratch glass, and have a metallic lustre. The
crystalline form of chromic oxide is identical with that of the
oxide of iron and alumina, with which it is isomorphous.

[7 bis] The most important of the compounds corresponding with chromic
oxide is _chromic chloride_, Cr_{2}Cl_{6}, which is known in an
anhydrous and in a hydrated form. It resembles ferric and aluminic
chlorides in many respects. There is a great difference between the
anhydrous and the hydrated chlorides; the former is insoluble in
water, the latter easily dissolves, and on evaporation its solution
forms a hygroscopic mass which is very unstable and easily evolves
hydrochloric acid when heated with water. The anhydrous form is of
a violet colour, and Wöhler gives the following method for its
preparation: an intimate mixture is prepared of the anhydrous
chromic oxide with carbon and organic matter, and charged into a
wide infusible glass or porcelain tube which is heated in a
combustion furnace; one extremity of the tube communicates with an
apparatus generating chlorine which is passed through several
bottles containing sulphuric acid in order to dry it perfectly
before it reaches the tube. On heating the portion of the tube in
which the mixture is placed and passing chlorine through, a
slightly volatile sublimate of chromic chloride, CrCl_{3} or
Cr_{2}Cl_{6}, is formed. This substance forms _violet tabular
crystals_, which may be distilled in dry chlorine without change,
but which, however, require a red heat for their volatilisation.
These crystals are greasy to the touch and insoluble in water, but
if they be powdered and boiled in water for a long time they pass
into _a green solution_. Strong sulphuric acid does not act on the
anhydrous salt, or only acts with exceeding slowness, like water.
Even aqua regia and other acids do not act on the crystals, and
alkalis only show a very feeble action. The specific gravity of the
crystals is 2·99. When fused with sodium carbonate and nitre they
give sodium chloride and potassium chromate, and when ignited in
air they form green chromic oxide and evolve chlorine. On ignition
in a stream of ammonia, chromic chloride forms sal-ammoniac and
chromium nitride, CrN (analogous to the nitrides BN, AlN). Mosberg
and Peligot showed that when chromic chloride is ignited in
hydrogen, it parts with one-third of its chlorine, forming chromous
chloride, CrCl_{2}--that is, there is formed from a compound
corresponding with chromic oxide, Cr_{2}O_{3}, a compound answering
to the _suboxide_, chromous oxide, CrO--just as hydrogen converts
ferric chloride into ferrous chloride with the aid of heat.
_Chromous chloride_, CrCl_{2}, forms colourless crystals easily
soluble in water, which in dissolving evolve a considerable amount
of heat, and form a blue liquid, capable of absorbing oxygen from
the air with great facility, being converted thereby into a chromic
compound.

The blue solution of chromous chloride may also be obtained by the
action of metallic zinc on the green solution of the hydrated
chromic chloride; the zinc in this case takes up chlorine just as
the hydrogen did. It must be employed in a large excess. Chromic
oxide is also formed in the action of zinc on chromic chloride, and
if the solution remain for a long time in contact with the zinc the
whole of the chromium is converted into chromic oxychloride. Other
chromic salts are also reduced by zinc into _chromous salts_,
CrX_{2}, just as the ferric salts FeX_{3} are converted into
ferrous salts FeX_{2} by it. The chromous salts are exceedingly
unstable and easily oxidise and pass into chromic salts; hence the
reducing power of these salts is very great. From cupric salts they
separate cuprous salts, from stannous salts they precipitate
metallic tin, they reduce mercuric salts into mercurous and ferric
into ferrous salts. Moreover, they absorb oxygen from the air
directly. With potassium chromate they give a brown precipitate of
chromium dioxide or of chromic oxide, according to the relative
amounts of the substances taken: CrO_{3} + CrO = 2CrO_{2} or
CrO_{3} + 3CrO = 2Cr_{2}O_{3}. Aqueous ammonia gives a blue
precipitate, and in the presence of ammoniacal salts a blue liquid
is obtained which turns red in the air from oxidation. This is
accompanied by the formation of compounds analogous to those given
by cobalt (Chapter XXII.). A solution of chromous chloride with a
hot saturated solution of sodium acetate, C_{2}H_{3}NaO_{2}, gives,
on cooling, transparent red crystals of chromous acetate,
C_{4}H_{6}CrO_{4},H_{2}O. This salt is also a powerful reducing
agent, but may be kept for a long time in a vessel full of carbonic
anhydride.

The insoluble anhydrous _chromic chloride_ CrCl_{3} very easily
_passes into solution_ in the presence of a trace (0·004) of
_chromous chloride_ CrCl_{2}. This remarkable phenomenon was
observed by Peligot and explained by Löwel in the following manner:
chromous chloride, as a lower stage of oxidation, is capable of
absorbing both oxygen and chlorine, combining with various
substances. It is able to decompose many chlorides by taking up
chlorine from them; thus it precipitates mercurous chloride from a
solution of mercuric chloride, and in so doing passes into chromic
chloride: 2CrCl_{2} + 2HgCl_{2} = Cr_{2}Cl_{6} + 2HgCl. Let us
suppose that the same phenomenon takes place when the anhydrous
chromic chloride is mixed with a solution of chromous chloride. The
latter will then take up a portion of the chlorine of the former,
and pass into a soluble hydrate of chromic chloride (hydrochloride
of oxide of chromium), and the original anhydrous chromic chloride
will pass into chromous chloride. The chromous chloride re-formed
in this manner will then act on a fresh quantity of the chromic
chloride, and in this manner transfer it entirely into solution as
hydrate. This view is confirmed by the fact that other chlorides,
capable of absorbing chlorine like chromous chloride, also induce
the solution of the insoluble chromic chloride--for example,
ferrous chloride, FeCl_{2}, and cuprous chloride. The presence of
zinc also aids the solution of chromic chloride, owing to its
converting a portion of it into chromous chloride. The solution of
chromic chloride in water obtained by these methods is perfectly
identical with that which is formed by dissolving chromic hydroxide
in hydrochloric acid. On evaporating the _green solution_ obtained
in this manner, it gives a green mass, containing water. On further
heating it leaves a soluble chromic oxychloride, and when ignited
it first forms an insoluble oxychloride and then chromic oxide; but
no anhydrous chromic chloride, Cr_{2}Cl_{6}, is formed by heating
the aqueous solution of chromic chloride, which forms an important
fact in support of the view that the green solution of chromic
chloride is nothing else but hydrochloride of oxide of chromium. At
100° the composition of the green hydrate is Cr_{2}Cl_{6},9H_{2}O,
and on evaporation at the ordinary temperature over H_{2}SO_{4}
crystals are obtained with 12 equivalents of water; the red mass
obtained at 120° contains Cr_{2}O_{3},4Cr_{2}Cl_{6},24H_{2}O. The
greater portion of it is soluble in water, like the mass which is
formed at 150°. The latter contains
Cr_{2}O_{3},2Cr_{2}Cl_{6},9H_{2}O = 3(Cr_{2}OCl_{4},3H_{2}O)--that
is, it presents the same composition as chromic chloride in which
one atom of oxygen replaces two of chlorine. And if the hydrate of
chromic chloride be regarded as Cr_{2}O_{3},6HCl, the substance
which is obtained should be regarded as Cr_{2}O_{3},4HCl combined
with water, H_{2}O. The addition of alkalis--for example,
baryta--to a solution of chromic chloride immediately produces a
precipitate, which, however, re-dissolves on shaking, owing to the
formation of one of the oxychlorides just mentioned, which may be
regarded as _basic salts_. Thus we may represent the product of the
change produced on chromic chloride under the influence of water
and heat by the following formulæ: first Cr_{2}O_{3},6HCl or
Cr_{2}Cl_{6},3H_{2}O is formed, then Cr_{2}O_{3},4HCl,H_{2}O or
Cr_{2}OCl_{4},3H_{2}O, and lastly Cr_{2}O_{3},2HCl,2H_{2}O or
Cr_{2}O_{2}Cl_{2},3H_{2}O. In all three cases there are 2
equivalents of chromium to at least 3 equivalents of water. These
compounds may be regarded as being intermediate between chromic
hydroxide and chloride; chromic chloride is Cr_{2}Cl_{6}, the first
oxychloride Cr_{2}(OH)_{2}Cl_{4}, the second Cr_{2}(OH)_{4}Cl_{2},
and the hydrate Cr_{2}(OH)_{6}--that is, the chlorine is replaced
by hydroxyl.

It is very important to remark two circumstances in respect to
this: (1) That the whole of the chlorine in the above compounds is
not precipitated from their solutions by silver nitrate; thus the
normal salt of the composition Cr_{2}Cl_{6},9H_{2}O only gives up
two-thirds of its chlorine; therefore Peligot supposes that the
normal salt contains the oxychloride combined with hydrochloric
acid: Cr_{2}Cl_{6} + 2H_{2}O = Cr_{2}O_{2}Cl_{2},4HCl, and that the
chlorine held as hydrochloric acid reacts with the silver, whilst
that held in the oxychloride does not enter into reaction, just as
we observe a very feebly-developed faculty for reaction in the
anhydrous chromic chloride; and (2) if the green aqueous solution
of CrCl_{3} be left to stand for some time, it ultimately turns
violet; in this form the whole of the chlorine is precipitated by
AgNO_{3}, whilst boiling re-converts it into the green variety.
Löwel obtained the violet solution of hydrochloride of chromic
oxide by decomposing the violet chromic sulphate with barium
chloride. Silver nitrate precipitates all the chlorine from this
violet modification; but if the violet solution be boiled and so
converted into the green modification, silver nitrate then only
precipitates a portion of the chlorine.

Recoura (1890-1893) obtained a crystallohydrate of violet chromium
sulphate, Cr_{2}(SO_{4})_{3}, with 18 or 15 H_{2}O. By boiling a
solution of this crystallohydrate, he converted it into the green
salt, which, when treated with alkalis, gave a precipitate of
Cr_{2}O_{3},2H_{2}O, soluble in 2H_{2}SO_{4} (and not 3), and only
forming the basic salt, Cr_{2}(OH)_{2}(SO_{4})_{2}. He therefore
concludes that the green salts are basic salts. The cryoscopic
determinations made by A. Speransky (1892) and Marchetti (1892)
give a greater 'depression' for the violet than the green salts,
that is, indicate a greater molecular weight for the green salts.
But as Étard, by heating the violet sulphate to 100°, converted it
into a green salt of the same composition, but with a smaller
amount of H_{2}O, it follows that the formation of a basic salt
alone is insufficient to explain the difference between the green
and violet varieties, and this is also shown by the fact that
BaCl_{2} precipitates the whole of the sulphuric acid of the violet
salt, and only a portion of that of the green salt. A. Speransky
also showed that the molecular electro-conductivity of the green
solutions is less than that of the violet. It is also known that
the passage of the former into the latter is accompanied by an
increase of volume, and, according to Recoura, by an evolution of
heat also.

Piccini's researches (1894) throw an important light upon the
peculiarities of the green chromium trichloride (or chromic
chloride); he showed (1) that AgF (in contradistinction to the
other salts of silver) precipitates all the chlorine from an
aqueous solution of the green variety; (2) that solutions of green
CrCl_{3},6H_{2}O in ethyl alcohol and acetone precipitate all their
chlorine when mixed with a similar solution of AgNO_{3}; (3) that
the rise of the boiling-point of the ethyl alcohol and acetone
green solutions of CrCl_{3},6H_{2}O (Chapter VII., Note 27 bis)
shows that i in this case (as in the aqueous solutions of MgSO_{4}
and HgCl_{2}) is nearly equal to 1, that is, that they are like
solutions of non-conductors; (4) that a solution of green CrCl_{3}
in methyl alcohol at first precipitates about 7/8 of its chlorine
(an aqueous solution about 2/3) when treated with AgNO_{3}, but
after a time the whole of the chlorine is precipitated; and (5)
that an aqueous solution of the green variety gradually passes into
the violet, while a methyl alcoholic solution preserves its green
colour, both of itself and also after the whole of the chlorine has
been precipitated by AgNO_{3}. If, however, in an aqueous or methyl
alcoholic solution only a portion of the chlorine be precipitated,
the solution gradually turns violet. In my opinion the general
meaning of all these observations requires further elucidation and
explanation, which should be in harmony with the theory of
solutions. Recoura, moreover, obtained compounds of the green salt,
Cr_{2}(SO_{4})_{3}, with 1, 2, and 3 molecules of H_{2}SO_{4},
K_{2}SO_{4}, and even a compound Cr_{2}(SO_{4})_{3}H_{2}CrO_{4}. By
neutralising the sulphuric acid of the compounds of
Cr_{2}(SO_{4})_{3} and H_{2}SO_{4} with caustic soda, Recoura
obtained an evolution of 33 thousand calories per each 2NaHO, while
free H_{2}SO_{4} only gives 30·8 thousand calories. Recoura is of
opinion that special _chromo sulphuric acids_, for instance
(CrSO_{4})H_{2}SO_{4} = 1/2Cr_{2}(SO_{4})_{3}H_{2}SO_{4}, are
formed. With a still larger excess of sulphuric acid, Recoura
obtained salts containing a still greater number of sulphuric acid
radicles, but even this method does not explain the difference
between the green and violet salts.

These facts must naturally be taken into consideration in order to
arrive at any complete decision as to the cause of the different
modifications of the chromic salts. We may observe that the green
modification of chromic chloride does not give double salts with
the metallic chlorides, whilst the violet variety forms compounds
Cr_{2}Cl_{6},2RCl (where R = an alkali metal), which are obtained
by heating the chromates with an excess of hydrochloric acid and
evaporating the solution until it acquires a violet colour. As the
result of all the existing researches on the green and violet
chromic salts, it appears to me most probable that their difference
is determined by the feeble basic character of chromic oxide, by
its faculty of giving basic salts, and by the colloidal properties
of its hydroxide (these three properties are mutually connected),
and moreover, it seems to me that the relation between the green
and violet salts of chromic oxide best answers to the relation of
the purpureo to the luteo cobaltic salts (Chapter XXII., Note 35).
This subject cannot yet be considered as exhausted (_see_ Note 7).

We may here observe that with tin the chromic salts, CrX_{3}, give
at low temperatures CrX_{2} and SnX_{2}, whilst at high
temperatures, on the contrary, CrX_{2} reduces the metal from its
salts SnX_{2}. The reaction, therefore, belongs to the class of
reversible reactions (Beketoff).

Poulenc obtained anhydrous CrF_{3} (sp. gr. 3·78) and CrF_{2} (sp.
gr. 4·11) by the action of gaseous HF upon CrCl_{2}. A solution of
fluoride of chromium is employed as a mordant in dyeing. Recoura
(1890) obtained green and violet varieties of Cr_{2}Br_{6},6H_{2}O.
The green variety can only be kept in the presence of an excess of
HBr in the solution; if alone its solution easily passes into the
violet variety with evolution of heat.

_Chromic oxide_, Cr_{2}O_{3}, rarely found, and in small quantities, in chrome ochre, is formed by the oxidation of chromium and its lower oxides, by the reduction of chromates (for example, of ammonium or mercuric chromate) and by the decomposition (splitting up) of the saline compounds of the oxide itself, CrX_{3} or Cr_{2}X_{6}, like alumina, which it resembles in forming a feeble base easily giving double and basic salts, which are either green or violet.

The reduction of chromic oxide--for instance, in a solution by zinc and sulphuric acid--leads to the formation of chromous oxide, CrO, and its salts, CrX_{2}, of a blue colour (_see_ Notes 7 and 7 ^{bis}). The further reduction[8] of oxide of chromium and its corresponding compounds gives _metallic chromium_. Deville obtained it (probably containing carbon) by reducing chromic oxide with carbon, at a temperature near the melting point of platinum, about 1750°, but the metal itself does not fuse at this temperature. Chromium has a steel-grey colour and is very hard (sp. gr. 5·9), takes a good polish, and dissolves in hydrochloric acid, but cold dilute sulphuric and nitric acids have no action upon it. Bunsen obtained metallic chromium by decomposing a solution of chromic chloride, Cr_{2}Cl_{6}, by a galvanic current, as scales of a grey colour (sp. gr. 7·3). Wöhler obtained crystalline chromium by igniting a mixture of the anhydrous chromic chloride Cr_{2}Cl_{6} (_see_ Note 7 bis) with finely-divided zinc, and sodium and potassium chlorides, at the boiling-point of zinc. When the resultant mass has cooled the zinc may be dissolved in dilute nitric acid, and grey crystalline chromium (sp. gr. 6·81) is left behind. Frémy also prepared crystalline chromium by the action of the vapour of sodium on anhydrous chromic chloride in a stream of hydrogen, using the apparatus shown in the accompanying drawing, and placing the sodium and the chromic chloride in separate porcelain boats. The tube containing these boats is only heated when it is quite full of dry hydrogen. The crystals of metallic chromium obtained in the tube are grey cubes having a considerable hardness and withstanding the action of powerful acids, and even of aqua regia. The chromium obtained by Wöhler by the action of a galvanic current is, on the contrary, acted on under these conditions. The reason of this difference must be looked for in the presence of impurities, and in the crystalline structure. But in any case, among the properties of metallic chromium, the following may be considered established: it is white in colour, with a specific gravity of about 6·7, is extremely hard in a crystalline form, is not oxidised by air at the ordinary temperature, and with carbon it forms alloys like cast iron and steel.

[8] The reduction of metallic chromium proceeds with comparative ease
in aqueous solutions. Thus the action of sodium amalgams upon a
strong solution of Cr_{2}Cl_{6} gives (first CrCl_{2}) an amalgam
of chromium from which the mercury may be easily driven off by
heating (in hydrogen to avoid oxidation), and there remains a
spongy mass of easily oxidizable chromium. Plaset (1891), by
passing an electric current through a solution of chrome alum mixed
with a small amount of H_{2}SO_{4} and K_{2}SO_{4}, obtained hard
scales of chromium of a bluish-white colour possessing great
hardness and stability (under the action of water, air, and acids).
Glatzel (1890) reduced a mixture of 2KCl + Cr_{2}Cl_{6} by heating
it to redness with shavings of magnesium. The metallic chromium
thus obtained has the appearance of a fine light-grey powder which
is seen to be crystalline under the microscope; its sp. gr. at 16°
is 6·7284. It fuses (with anhydrous borax) only at the highest
temperatures, and after fusion presents a silver-white fracture.
The strongest magnet has no action upon it.

Moissan (1893) obtained chromium by reducing the oxide Cr_{2}O_{3}
with carbon in the electrical furnace (Chapter VIII., Note 17) in
9-10 minutes with a current of 350 ampères and 50 volts. The
mixture of oxide and carbon gives a bright ingot weighing 100-110
grams. A current of 100 ampères and 50 volts completes the
experiment upon a smaller quantity of material in 15 minutes; a
current of 30 ampères and 50 volts gave an ingot of 10 grams in
30-40 minutes. The resultant carbon alloy is more or less rich in
chromium (from 87·37-91·7 p.c.). To obtain the metal free from
carbon, the alloy is broken into large lumps, mixed with oxide of
chromium, put into a crucible and covered with a layer of oxide.
This mixture is then heated in the electric furnace and the pure
metal is obtained. This reduction can also be carried on with
chrome iron ore FeOCr_{2}O_{3} which occurs in nature. In this case
a homogeneous alloy of iron and chromium is obtained. If this alloy
be thrown in lumps into molten nitre, it forms insoluble
sesquioxide of iron and a soluble alkaline chromate. This alloy of
iron and chromium dissolved in molten steel (chrome steel) renders
it hard and tough, so that such steel has many valuable
applications. The alloy, containing about 3 p.c. Cr and about 1·3
p.c. carbon, is even harder than the ordinary kinds of tempered
steel and has a fine granular fracture. The usual mode of preparing
the ferrochromes for adding to steel is by fusing powdered chrome
iron ore under fluxes in a graphite crucible.

The two analogues of chromium, _molybdenum_ and _tungsten_ (or wolfram), are of still rarer occurrence in nature, and form acid oxides, RO_{3}, which are still less energetic than CrO_{3}. Tungsten occurs in the somewhat rare minerals, _scheelite_, CaWO_{4}, and _wolfram_; the latter being an isomorphous mixture of the normal tungstates of iron and manganese, (MnFe)WO_{4}. Molybdenum is most frequently met with as _molybdenite_, MoS_{2}, which presents a certain resemblance to graphite in its physical properties and softness. It also occurs, but much more rarely, as a yellow lead ore, PbMoO_{4}. In both these forms molybdenum occurs in the primary rocks, in granites, gneiss, &c., and in iron and copper ores in Saxony, Sweden, and Finland. Tungsten ores are sometimes met with in considerable masses in the primary rocks of Bohemia and Saxony, and also in England, America, and the Urals. The preliminary treatment of the ore is very simple; for example, the sulphide, MoS_{2}, is roasted, and thus converted into sulphurous anhydride and molybdic anhydride, MoO_{3}, which is then dissolved in alkalis, generally in ammonia. The ammonium molybdate is then treated with acids, when the sparingly soluble molybdic acid is precipitated. Wolfram is treated in a different manner. Most frequently the finely-ground ore is repeatedly boiled with hydrochloric and nitric acids, and the resultant solutions (of salts of manganese and iron) poured off, until the dark brown mass of ore disappears, whilst the tungstic acid remains, mixed with silica, as an insoluble residue; it is treated also with ammonia, and is thus converted into soluble ammonium tungstate, which passes into solution and yields tungstic acid when treated with acids. This hydrate is then ignited, and leaves tungstic anhydride. The general character of molybdic and tungstic anhydrides is analogous to that of chromic anhydride; they are anhydrides of a feebly acid character, which easily give polyacid salts and colloid solutions.[8 bis]

[8 bis] The atomic composition of the tungsten and molybdenum compounds
is taken as being identical with that of the compounds of sulphur
and chromium, because (1) both these metals give two oxides in
which the amounts of oxygen per given amount of metal stand in the
ratio 2 : 3; (2) the higher oxide is of the latter kind, and, like
chromic and sulphuric anhydrides, it has an acid character; (3)
certain of the molybdates are isomorphous with the sulphates; (4)
the specific heat of tungsten is 0·0334, consequently the product
of the atomic weight and specific heat is 6·15, like that of the
other elements--it is the same with molybdenum, 96·0 × 0·0722 =
6·9; (5) tungsten forms with chlorine not only compounds WCl_{6},
WCl_{5}, and WOCl_{4}, but also WO_{2}Cl_{2}, a volatile substance
the analogue of chromyl chloride, CrO_{2}Cl_{2}, and sulphuryl
chloride, SO_{2}Cl_{2}. Molybdenum gives the chlorine compounds,
MOCl_{2}, MOCl_{3}(?), MOCl_{4} (fuses at 194°, boils at 268°;
according to Debray it contains MOCl_{5}), MoOCl_{4},
MoO_{2}Cl_{2}, and MoO_{2}(OH)Cl. The existence of tungsten
hexachloride, WCl_{6}, is an excellent proof of the fact that the
type SX_{6} appears in the analogues of sulphur as in SO_{3}; (6)
the vapour density accurately determined for the chlorine compounds
MoCl_{4}, WCl_{6}, WCl_{5}, WOCl_{4} (Roscoe) leaves no doubt as to
the molecular composition of the compounds of tungsten and
molybdenum, for the observed and calculated results entirely agree.

Tungsten is sometimes called scheele in honour of Scheele, who
discovered it in 1781 and molybdenum in 1778. Tungsten is also
known as wolfram; the former name was the name given to it by
Scheele, because he extracted it from the mineral then known as
tungsten and now called scheelite, CaWO_{4}. The researches of
Roscoe, Blomstrand and others have subsequently thrown considerable
light on the whole history of the compounds of molybdenum and
tungsten.

The ammonium salts of tungsten and molybdic acids when ignited
leave the anhydrides, which resemble each other in many respects.
_Tungsten anhydride_, WO_{3}, is a yellowish substance, which only
fuses at a strong heat, and has a sp. gr. of 6·8. It is insoluble
both in water and acid, but solutions of the alkalis, and even of
the alkali carbonates, dissolve it, especially when heated, forming
alkaline salts. _Molybdic anhydride_, MoO_{3}, is obtained by
igniting the acid (hydrate) or the ammonium salt, and forms a white
mass which fuses at a red heat, and solidifies to a yellow
crystalline mass of sp. gr. 4·4; whilst on further heating in open
vessels or in a stream of air this anhydride _sublimes_ in pearly
scales--this enables it to be obtained in a tolerably pure state.
Water dissolves it in small quantities--namely, 1 part requires 600
parts of water for its solution. The hydrates of molybdic anhydride
are _soluble also in acids_ (a hydrate, H_{2}MoO_{4}, is obtained
from the nitric acid solution of the ammonium salt), which forms
one of their distinctions from the tungstic acids. But after
ignition molybdic anhydride is insoluble in acids, like tungstic
anhydride; alkalis dissolve this anhydride, easily forming
molybdates. Potassium bitartrate dissolves the anhydride with the
aid of heat. None of the acids yet considered by us form so many
different salts with one and the same base (alkali) as molybdic and
tungstic acids. The composition of these salts, and their
properties also, vary considerably. The most important discovery in
this respect was made by Marguerite and Laurent, who showed that
the salts which contain a large proportion of tungstic acid are
easily soluble in water, and ascribed this property to the fact
that tungstic acid may be obtained _in several states_. The common
tungstates, obtained with an excess of alkali, have an alkaline
reaction, and on the addition of sulphuric or hydrochloric acid
first deposit an acid salt and then a hydrate of tungstic acid,
which is insoluble both in water and acids; but if instead of
sulphuric or hydrochloric acids, we add acetic or phosphoric acid,
or if the tungstate be saturated with a fresh quantity of tungstic
acid, which may be done by boiling the solution of the alkali salt
with the precipitated tungstic acid, a solution is obtained which,
on the addition of sulphuric or a similar acid, does not give a
precipitate of tungstic acid at the ordinary or at higher
temperatures. The solution then contains peculiar salts of tungstic
acid, and if there be an excess of acid it also contains tungstic
acid itself; Laurent, Riche, and others called it _metatungstic
acid_, and it is still known by this name. Those salts which with
acids immediately give the insoluble tungstic acid have the
composition R_{2}WO_{4},RHWO_{4}, whilst those which give the
soluble metatungstic acid contain a far greater proportion of the
acid elements. Scheibler obtained the (soluble) metatungstic acid
itself by treating the soluble barium (meta) tetratungstate,
BaO,4WO_{3}, with sulphuric acid. Subsequent research showed the
existence of a similar phenomenon for molybdic acid. There is no
doubt that this is a case of colloidal modifications.

Many chemists have worked on the various salts formed by molybdic
and tungstic acids. The tungstates have been investigated by
Marguerite, Laurent, Marignac, Riche, Scheibler, Anthon, and
others. The molybdates were partially studied by the same chemists,
but chiefly by Struvé and Svanberg, Delafontaine, and others. It
appears that for a given amount of base the salts contain one to
eight equivalents of molybdic or tungstic anhydride; _i.e._ if the
base have the composition RO, then the highest proportion of base
will be contained by the salts of the composition ROWO_{3} or
ROMoO_{3}--that is, by those salts which correspond with the normal
acids H_{2}WO_{4} and H_{2}MoO_{4}, of the same nature as sulphuric
acid; but there also exist salts of the composition RO,2WO_{3},
RO,3WO_{3} ... RO,8WO_{3}. The water contained in the composition
of many of the acid salts is often not taken into account in the
above. The properties of the salts holding different proportions of
acids vary considerably, but one salt may be converted into another
by the addition of acid or base with great facility, and the
greater the proportion of the elements of the acid in a salt, the
more stable, within a certain limit, is its solution and the salt
itself.

The most common ammonium molybdate has the composition
(NH_{4}HO)_{6},H_{2}O,7MoO_{3} (or, according to Marignac and
others, NH_{4}HMoO_{4}), and is prepared by evaporating an
ammoniacal solution of molybdic acid. It is used in the laboratory
for precipitating phosphoric acid, and is purified for this purpose
by mixing its solution with a small quantity of magnesium nitrate,
in order to precipitate any phosphoric acid present, filtering, and
then adding nitric acid and evaporating to dryness. A pure ammonium
molybdate free from phosphoric acid may then be extracted from the
residue.

Phosphoric acid forms insoluble compounds with the oxides of
uranium and iron, tin, bismuth, &c., having feeble basic and even
acid properties. This perhaps depends on the fact that the atoms of
hydrogen in phosphoric acid are of a very different character, as
we saw above. Those atoms of hydrogen which are replaced with
difficulty by ammonium, sodium, &c., are probably easily replaced
by feebly energetic acid groups--that is, the formation of
particular complex substances may be expected to take place at the
expense of these atoms of the hydrogen of phosphoric acid and of
certain feeble metallic acids; and these substances will still be
acids, because the hydrogen of the phosphoric acids and metallic
acids, which is easily replaced by metals, is not removed by their
mutual combination, but remains in the resultant compound. Such a
conclusion is verified in the _phosphomolybdic acids_ obtained
(1888) by Debray. If a solution of ammonium molybdate be acidified,
and a small amount of a solution (it may be acid) containing
orthophosphoric acid or its salts be added to it (so that there are
at least 40 parts of molybdic acid present to 1 part of phosphoric
acid), then after a period of twenty-four hours the whole of the
phosphoric acid is separated as a yellow precipitate, containing,
however, not more than 3 to 4 p.c. of phosphoric anhydride, about 3
p.c. of ammonia, about 90 p.c. of molybdic anhydride, and about 4
p.c. of water. The formation of this precipitate is so distinct and
so complete that this method is employed for the discovery and
separation of the smallest quantities of phosphoric acid.
Phosphoric acid was found to be present in the majority of rocks by
this means. The precipitate is soluble in ammonia and its salts, in
alkalis and phosphates, but is perfectly insoluble in nitric,
sulphuric, and hydrochloric acids in the presence of ammonium
molybdate. The composition of the precipitate appears to vary under
the conditions of its precipitation, but its nature became clear
when the acid corresponding with it was obtained. If the
above-described yellow precipitate be boiled in aqua regia, the
ammonia is destroyed, and an acid is obtained in solution, which,
when evaporated in the air, crystallises out in yellow oblique
prisms of approximately the composition
P_{2}O_{5},20MoO_{3},26H_{2}O. Such an unusual proportion of
component parts is explained by the above-mentioned considerations.
We saw above that molybdic acid easily gives salts
R_{2}O_n_MoO_{3}_m_H_{2}O, which we may imagine to correspond to a
hydrate MoO_{2}(HO)_{2}_n_Mo_{3}_m_H_{2}O. And suppose that such a
hydrate reacts on orthophosphoric acid, forming water and compounds
of the composition MoO_{2}(HPO_{4})_n_MoO_{3}_m_H_{2}O or
MoO_{3}(H_{2}PO_{4})_{2}_n_MoO_{3}_m_H_{2}O; this is actually the
composition of phosphomolybdic acid. Probably it contains a portion
of the hydrogen replaceable by metals of both the acids H_{3}PO_{4}
and of H_{2}MoO_{4}. The crystalline acid above is probably
H_{3}MoPO_{7},9MoO_{3},12H_{2}O. This acid is really tribasic,
because its aqueous solution precipitates salts of potassium,
ammonium, rubidium (but not lithium and sodium) _from acid
solutions_, and gives a _yellow_ precipitate of the composition
R_{3}MoPO_{7},9MoO_{3},8H_{2}O, where R = NH_{4}. Besides these,
salts of another composition may be obtained, as would be expected
from the preceding. These salts are only stable in acid solutions
(which is naturally due to their containing an excess of acid
oxides), whilst under the action of alkalis they give _colourless_
phosphomolybdates of the composition R_{3}MoPO_{3},MoO_{2},3H_{2}O.
The corresponding salts of potassium, silver, ammonium, are easily
soluble in water and crystalline.

Phosphomolybdic acid is an example of the _complex inorganic acids_
first obtained by Marignac and afterwards generalised and studied
in detail by Gibbs. We shall afterwards meet with several examples
of such acids, and we will now turn attention to the fact that they
are usually formed by weak polybasic acids (boric, silicic,
molybdic, &c.), and in certain respects resemble the cobaltic and
such similar complex compounds, with which we shall become
acquainted in the following chapter. As an example we will here
mention certain complex compounds containing molybdic and tungstic
acids, as they will illustrate the possibility of a considerable
complexity in the composition of salts. The action of ammonium
molybdate upon a dilute solution of purpureocobaltic salts (_see_